Shielding, Effective Nuclear Charge, and Periodic Trends Flashcards
What are the important factors when atoms come together?
atomic size, ionization energy (ease of losing an electron, how tightly the valence electrons are held), and electron affinity (depends on effective nuclear charge and distance from nucleus)
What is shielding?
Core electrons block valence electrons from experiencing the full attraction of the nucleus, they cancel some of the positive nuclear charge.
What is effective nuclear charge?
Also known as Z*eff. It is the actual positive nuclear charge felt by the electron in the atom (other than hydrogen).
How is effective nuclear charge calculated?
Z = atomic number/nuclear charge
S = number of shielding electrons
Z*eff = Z - S
Do electrons in the same shell shield each other?
No (valence electrons don’t contribute to shielding either)
How is atomic radius measured if the electron cloud around the nucleus has no definite limit/edge?
Atomic radius is half the distance between two identical nuclei solids (metallic). For non-metallic it is the distance from the middle of both to halfway of one blob.
What factors affect atomic size?
- Principal quantum number for valence electrons
- effective nuclear charge felt by valence electrons
How does principal quantum number affect atomic size?
As n increases or decreases, so does the atomic orbital.
- orbitals are bigger towards the bottom of the periodic table because the number of electron shells increase and valence electrons are farther away (larger atomic radius)
- As you go left to right, atomic size gets smaller because there are more protons (nuclear charge) and therefore and so electrons are pulled more tightly to the nucleus (smaller radius)
How does effective nuclear charge felt by valence electrons affect atomic size?
- From top to bottom nuclear charge stays mostly the same
- From left to right, nuclear charge gets bigger so there are more protons and electrons are held closer (smaller radius). Nuclear charge increases as core electrons stay the same
Why are cations smaller than the original atoms?
The number of protons exceeds the number of electrons, since there are less of the electrons they are pulled closer to the nucleus (smaller radius). Size decreases as you go right and increases as you go down
Why are anions larger than the original atoms?
Number of protons is less than the number of electrons so the electrons are less pulled towards the nucleus as well as more repulsion between the electrons (larger radius). Anions increase in size as you go down
What is ionization energy? X (g) –> X (g) + é
The energy required to remove electrons from isolated gaseous atoms/ions.
- Written in kj/mol (change in heat when an electron is removed)
- Reflects how tightly an electron is held by the nucleus
- Endothermic (positive value)
- Increases left to right as nuclear charge increases
- Decreases top to bottom as n increases (larger atom smaller IE)
What is the first ionization energy?
Energy to remove the most loosely bound electron from a gaseous atom in ground state.
- Successive IE are possible until no electron remains, you also need more energy to remove more and more electrons
What elements have the largest ionization energy?
- Non-metals have the largest IE, metals have the smallest IE
- Trends in IE are opposite the atomic size
When does it require more energy to move an electron?
From a positive ion and core electrons
What is electron affinity? X (g) + é –> X (g)
The desire of a gaseous atom/ion for adding an electron to its valence shell
- Measured in kj/mol (heat released/absorbed when one mole electron is added to gaseous atom/ion)
- For most elements it is exothermic, some are endothermic though (no desire)
- Increases left to right as nuclear charge increases
What elements have a higher affinity for electrons?
Non-metals have a higher E.A, metals don’t like to form anions
